91Ó°ÊÓ

In any chemical reaction, the limiting reactant is the substance that is completely consumed first, halting the reaction and determining the amount of product formed. To find the limiting reactant, it is crucial to compare the available moles of each reactant with the proportions required by the balanced chemical equation.

In our example with ammonia (\(\text{NH}_3\)) and oxygen (\(\text{O}_2\)), the balanced equation requires 4 moles of \(\text{NH}_3\) for every 5 moles of \(\text{O}_2\). With the given amounts, we calculated 0.587 moles of \(\text{NH}_3\) and 0.625 moles of \(\text{O}_2\).

• To fully react 0.587 moles of \(\text{NH}_3\), we'd need \(\frac{5}{4} \times 0.587 ≈ 0.734\) moles of \(\text{O}_2\).
• However, we only have 0.625 moles of \(\text{O}_2\), making \(\text{O}_2\) the limiting reactant.

The limiting reactant dictates the extent of the reaction because it runs out first, preventing more products from forming. Thus, in this scenario, the reaction will stop once all 0.625 moles of \(\text{O}_2\) is consumed.

Molar mass is a fundamental concept in chemistry that represents the mass of one mole of a substance, usually expressed in grams per mole (g/mol). Calculating the molar mass is the initial step in converting between the mass of a substance and moles, which are used to describe quantities in chemical reactions.

For example, the molar mass of ammonia (\(\text{NH}_3\)) can be found by summing the atomic masses of nitrogen and hydrogen:

• Nitrogen (\(\text{N}\)): 14.01 g/mol
• Hydrogen (\(\text{H}\)): 1.01 g/mol (each hydrogen atom)

For \(\text{NH}_3\), add the total for hydrogen (1.01 g/mol \(\times\) 3) and nitrogen, giving a molar mass of 17.03 g/mol.

Similarly, the molar mass of oxygen (\(\text{O}_2\)) is calculated by multiplying the atomic mass of one oxygen by two since \(\text{O}_2\) is a diatomic molecule:

• Oxygen: 16.00 g/mol × 2 = 32.00 g/mol

These calculations allow us to convert grams of ammonia and oxygen into moles, facilitating the determination of the limiting reactant and the progress of the chemical reaction.

An exothermic reaction is characterized by the release of energy, typically in the form of heat, into the surroundings. This type of reaction is indicated by a negative enthalpy change (\(\Delta H\)). In our exercise, the enthalpy change of \(-906 \text{ kJ}\) suggests that when ammonia reacts with oxygen, a significant amount of energy is released as heat.

Understanding the details of exothermic reactions can help predict the temperature changes and energy dynamics within the reaction:

• The negative sign of \(\Delta H\) (-906 \(\text{kJ}\)) indicates the reaction releases rather than absorbs energy.
• The magnitude of the energy change helps gauge how much heat will be expelled.
• The greater the exothermicity, the more significant the temperature rise in the reaction mixture and its surroundings.

The amount of heat released is calculated based on the limiting reactant, \(\text{O}_2\) in this case, which determines the usable portion of the reactants. For each mole of \(\text{O}_2\) consumed, the equation provides \(\frac{-906}{5} \text{ kJ} ≈ -181.2\text{ kJ/mole}\). Therefore, the reaction of 0.625 moles of \(\text{O}_2\) results in the release of approximately \(-113.3 \text{ kJ}\) of energy.